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Write formulas and names for ionic, polyatomic and covalent compounds using IUPAC nomenclature.

UNIT 1 Chemical Reactions Part I (Text: p. 276-327). Write formulas and names for ionic, polyatomic and covalent compounds using IUPAC nomenclature. Write and classify balanced chemical equations from written descriptions of reactions.

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Write formulas and names for ionic, polyatomic and covalent compounds using IUPAC nomenclature.

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  1. UNIT 1 Chemical Reactions Part I (Text: p. 276-327) • Write formulas and names for ionic, polyatomic and covalent compounds using IUPAC nomenclature. • Write and classify balanced chemical equations from written descriptions of reactions. • Predict the products of chemical reactions, given the reactants and type of reaction. • Determine the average atomic mass using isotopes and their relative abundance. • Research the importance and application of isotopes. • Describe the concept of the mole and its importance to measurement in chemistry. • Calculate the molar mass of various substances. • Solve problems requiring conversions between moles, mass, and number of particles.

  2. Anatomy of an Atom All atoms, except hydrogen, are made of 3 basic particles: protons, neutrons and electrons. Each element has a unique number of protons, which is indicated by its atomic number.

  3. The outermost shell is called the ________________. The electrons in the valence shell are called the _________________________ The atoms of elements in Period 1 have one shell. This shell contains a maximum of ____ electrons. The atoms of elements in Period 2 have _____ shells. The valence shell contains a maximum of ______ electrons. The atoms of elements in Period 3 have _____ shells. The valence shell contains a maximum of ______ electrons. valence shell valence electrons 2 2 8 3 8 *Complete Electron Shell handout*

  4. Chemical Formulas • Chemistry has its own language. Chemists communicate in this language to describe the millions of known compounds. This communication depends on a standard system of naming and writing the formulas for compounds. Chemists formed a group to standardize the system of naming and called themselves the International Union of Physical and Applied Chemists, or IUPAC. • A chemical formula is a shorthand method to represent compounds that uses the elements' symbols and subscripts. The chemical formula gives the following information: • The different elements in the compound. • The number of atoms of each element in the compound.

  5. element symbols • Subscript tells you amount of each element. • Water contains: • 2 Hydrogen’s • 1 Oxygen H2O subscript * No subscript indicates only 1 atom is present* Ca3(PO4)2 • Contains 3 calcium atoms *every subscript inside the brackets needs to be multiplied by 2* P 1 atom x 2 = 2 phosphorus atoms O 4 atoms x 2 = 8 oxygen atoms

  6. Ionic Compounds An __________ is a charged particle. An ion is formed when a neutral atom gains or loses electrons. Positively charged ions are called _____________, and negatively charged ions are called ________________. ion cations anions *Complete “Keeping an ION That” handout*

  7. Ionic compounds are formed when two or more oppositely charged ions are attracted to each other. This chemical attraction is called a chemical bond. An ionic bond is formed when a negatively charged ion is attracted to a positively charged ion. Ions combine together so that their charges add to zero. Ionic compounds are usually made of _______________________ ions. metal and non-metal Example: NaCl – sodium chloride Fe2O3 – iron oxide CuSO4 – copper sulfate Ca3(PO4)2 – calcium phosphate m nm m nm

  8. The periodic table arranges atoms according to their properties. The periodic table below shows the names of several groups we will be referring to throughout this course. metals non-metals

  9. Writing Binary Ionic Formula A binary compound contains ______ different kinds of elements. There can be more than one atom of each element in a binary compound. Binary ionic compounds usually contain one kind of metal ion combined with one kind of non-metal ion. Metal ions have ________________ charges and non-metal ions have ________________ charges. 2 positive negative • When naming an ionic compound from its formula follow the rules below: • The cation (positive ion) is named first, followed by the anion (negative ion). • Write the full name of the metallic element (positive ion). • Write the name of the non-metallic element (negative ion) and change the ending to "-ide".

  10. Example: Write the name of NaCl. Step 1: Name the first element. Step 2: Name the second element and change the ending to "-ide". The name of the compound is ________________________. Na = sodium Cl = chlorine  chloride sodium chloride Example: Write the name of Mg3P2. Step 1: Name the first element. Step 2: Name the root of the second element and add "-ide". The name of the compound is ____________________________. Mg = magnesium P = phosphorus  phosphide magnesium phosphide

  11. Naming Binary Ionic Formula There are two methods for determining the formula of a compound, but the following points must hold true: • The formula must have the cation first, followed by the anion. • The sum of the charges of the ions must be zero. That is, the number of positive charges must equal the number of negative charges. • You may not change the charge of the ions to make the ion charges equal zero. • Method 1 – Lowest Common Multiple • Write the symbols for the ions involved. • Determine the lowest whole number ratio that will give an overall net charge of zero. That is, the number of positive charges must equal the number of negative charges.

  12. Example: Write the formula for aluminum oxide. Step 1: Write the ions and their charges. Step 2: Determine the lowest common multiple. The formula for aluminum oxide is ___________________. Al3+ and O2- • LCM for 3 and 2 is 6 • To get 6 positive, need 2 Al3+ 2 x 3+ = 6+ • To get 6 negative, need 3 O2-  3 x 2- = 6- becomes subscript Al2O3

  13. Method 2 – The "Criss-Cross" Method This method accomplishes the same as the lowest common multiple method-the total charge of the compound is zero. • Write the ions and their charges side by side. • Make the number of the charge of one ion the subscript of the other ion (omitting the + or – sign). Remember we do not write the number one as a subscript. • Reduce all subscripts to their simplest form, if necessary.

  14. Example: Write the formula for aluminum oxide. Step 1: Write the ions and their charges. Step 2: Make the number of the charge of one ion the subscript of the other ion. Example: Write the formula for barium fluoride. Note: The charge on the fluoride ion is 1–. Since IUPAC rules do not write the number one as a subscript, we leave the barium without a subscript. Al3+ and O2- Al3+ O2- Al2O3 Ba2+ F1- BaF2

  15. Polyatomic Ions Some ions are composed of several atoms joined covalently. These are called polyatomic ions (poly = many). Refer to the Table of Common Polyatomic Ions for a list of ions. The charge for polyatomic ions is for the whole group of atoms not just for the atom written last. DO NOT change the subscripts of polyatomic ions; if you change the subscripts you change the identity of these ions. When indicating the presence of more than one polyatomic ion in a compound, we use parenthesis around the polyatomic ion, followed by its subscript. For example, the compound Al(C2H3O2)3 has an aluminum ion and 3 acetate ions. Placing the acetate ion in parenthesis and following it with the subscript 3 indicates there are 3 acetate ions.

  16. Example: Write the name for KNO3. Step 1: Identify the cation. Step 2: Identify the anion. Step 3: Write the name of the cation first, followed by the anion. K+  potassium ion NO3-  nitrate ion potassium nitrate Example: Write the name of Hg2Cl2. Step 1: Identify the cation. Step 2: Identify the anion. Step 3: Write the name of the cation first, followed by the anion. Hg22+ dimercury ion Cl- chloride ion dimercury chloride

  17. Example: Write the name of Na3PO4. Step 1: Identify the cation. Step 2: Identify the anion. Step 3: Write the name of the cation first, followed by the anion. Na+  sodium ion PO43-  phosphate ion sodium phosphate Example: Write the name of NH4SCN. Step 1: Identify the cation. Step 2: Identify the anion. Step 3: Write the name of the cation first, followed by the anion. NH4+  ammonium ion SCN-  thiocyanate ion ammonium thiocyanate *Complete Ionic Compound Worksheet*

  18. Stock Naming System Most of the transition metals have more than one possible ion charge. They are often referred to as being multivalent. For example, In 1919, Alfred Stock (1876 – 1946), a German chemist, suggested using numbers to indicate the charge of the ions. Prior to this the ions were given different names based upon their charge. The Cu+ ion was called cuprous and the Cu2+ ion was called cupric. However, the Fe2+ ion was ferrous and the Fe3+ ion was ferric. Since the charges were not always the same, the "–ic" and "–ous" suffixes caused some confusion. Today, the Stock naming system uses Roman numerals following the metal ion's name to indicate the ion's charge.

  19. Example: Copper (I) = Cu+ Copper (II) = Cu2+ Iron (II) = Fe2+Iron (III) = Fe3+ As a general rule, all metals are multivalent (have more than one ion charge) except group one and two metals, silver, cadmium, zinc, and aluminum. Unless the metal is one of these use the Roman numeral. cuprous cupric ferrous ferric

  20. Multivalent Compounds Example: Write the formula for iron (III) chloride Step 1: Write out the ions. Step 2: Balance or “criss-cross” the charges. Fe3+ and Cl- FeCl3 Example: Write the formula for lead (IV) sulfide. Step 1: Write the ions. Step 2: Balance or “criss-cross” the charges. Step 3: Reduce the subscripts. Pb4+ and S2- Pb2S4 *divide by GCF* Pb2S 2  PbS2

  21. Naming Multivalent Compounds • We name in a very similar manner as those ions with a single ion charge, except we must determine the charge on the metal ion. • To determine the charge on the metal ion, • Write the name of the ions. • Multiply the charge of the anion by its subscript. • Divide this number by the subscript of the metal ion. The result is the   • charge on the metal ion.

  22. Example: Write the name for CoBr2. Step 1: Write the names of the ions. (including the charge of the anion) Step 2: Multiply the charge of the bromide by its subscript then divide by the subscript for cobalt. Step 3: Write the name, indicating the charge of cobalt using roman numerals. cobalt bromide Cox Br1- 1 x 2 = 2 1 = 2 *ignore charge* charge for cobalt cobalt (II) bromide

  23. Example: Write the name for MnO2. Step 1: Write the names of the ions. Step 2: Multiply the charge of the oxide by its subscript then divide by the subscript for manganese. Step 3: Write the name. manganese oxide Mnx O2- 2 x 2 = 4 1 = 4 charge for manganese manganese (IV) oxide *complete Multivalent Worksheet*

  24. Naming Covalent Compounds Example: Write the name for CO2. Step 1: Name the first atom with prefixes. Step 2: Name the second element using prefixes and end in "-ide". Step 3: Write the name of the compound writing the substance found more to the left on the periodic table first. There is only 1 carbon. We omit “mono” for the first element. carbon There are 2 oxygen's, so we use the di prefix dioxide carbon dioxide

  25. Example: Write the name for N2O4. Step 1: Name the first atom with prefixes. Step 2: Name the second element using prefixes and end in "-ide". Step 3: Write the name of the compound. dinitrogen tetraoxide dinitrogen tetraoxide

  26. Writing Covalent Compound Formulas • Writing formulas for covalent compounds involves the following rules: • Write the symbol for the first element followed by the subscript indicated by the prefix. • Write the symbol of the second element followed by the subscript indicated by its prefix. DO NOT REDUCE THE SUBSCRIPTS!!! Example: Write the formula for dinitrogen monoxide. Step 1: Write the symbol and subscript for the first element. Step 2: Write the symbol and subscript for the second element. Step 3: Combine dinitrogen  N2 monoxide  O * 1 as a subscript is not needed * N2O Example: Write the formula for sulphur hexafluoride. SF6

  27. Diatomic Molecules Some elements do not exist as single atoms. These elements exist as pairs of atoms joined covalently, called diatomic molecules. The elements that exist as diatomic molecules are hydrogen (H2), oxygen (O2), fluorine (F2), chlorine (Cl2), bromine (Br2), iodine (I2) and nitrogen (N2). When oxygen gas, hydrogen gas, etc. is used the formula will be O2, H2, etc. Here is a mnemonic device to help you remember the diatomic molecules: IHave No Bright Or Clever Friends I = iodine H = hydrogen N = nitrogen Br = bromine O= oxygen Cl = chlorine F = fluorine * Complete Covalent Compound Worksheet *

  28. Chemical Equations A chemical equation indicates the substances reacting and the substances produced in a chemical reaction. A chemical equation also shows the ratio in which these substances react or are produced. A word equation can describe a chemical reaction. Example: Hydrogen gas and oxygen gas react to form (or yield) water vapour. A chemical equation can also show heat changes that occur. Endothermic reactions cause the reaction vessel to feel cooler because the reaction absorbs energy. The energy is used in the reaction, so energy is a reactant. Exothermic reactions release energy. Consequently, we consider heat or energy to be a product of an exothermic reaction. 2 H2(g) + O2(g)  2 H2O(g) 2 H2(g) + O2(g) 2 H2O(g) + energy

  29. Recall: and, reacts with, separates 2 or more reactants yields, forms, separates reactants and products solid state liquid state gaseous state aqueous state (in water)

  30. Balancing Chemical Equations Example: Balance the equation C3H8 + O2 CO2 + H2O * See white board* Example: Balance the equation Al2(SO4)3 + CaCl2  AlCl3 + CaSO4 * See white board* * Complete Balancing Chemical Equations Worksheet*

  31. Types of Chemical Reactions Combustion Reactions Are very rapid reactions of a hydrocarbon (fuel) substance with oxygen gas producing carbon dioxide and water plus a lot of heat. Example: CH4(g) (methane) + O2(g) CO2(g) + H2O(g) Synthesis Reactions Involve the combining of smaller atoms/molecules into larger, more complex molecules. If only two different atoms appear on the reactant side, then the reaction must be synthesis. Example: Water and dinitrogen pentoxide gas react to produce aqueous hydrogen nitrate. 2 2 H2O (l) + N2O5 (g)  2 HNO3 (aq)

  32. Decomposition Reactions Involve the splitting of large molecules into smaller molecules or elements. Example: Solid nickel (II) hydroxide decomposes to produce solid nickel (II) oxide and water. Single Displacement Reactions Are chemical changes that involve an element and a compound as reactants. Example: Fluorine gas will react with sodium bromide in an aqueous solution to produce sodium fluoride and liquid bromine. Ni(OH)2 (s)  NiO (s) + H2O (l) F2 (g) + 2 NaBr (aq)  2 NaF (aq) + Br2 (l) switch

  33. Double Displacement Reactions Occur when elements in different compounds displace each other or exchange places. Generally, the reaction occurs in an aqueous system. Example: When aqueous lithium iodide and aqueous silver nitrate react, they will produce solid silver iodide and aqueous lithium nitrate. Acid-Base Reactions When an acid and base react together, the reaction is known as a neutralization reaction. The products will always be water and a salt. Example: When a solution of aqueous hydrochloric acid and solid potassium hydroxide react, water and aqueous potassium chloride are formed. LiI (aq) + AgNO3 (aq)  AgI (s) + LiNO3 (aq) acid base water salt HCl (aq) + KOH (s)  H2O (l) + KCl (aq)

  34. Isotopes • The number of neutrons in each atom varies, even between atoms of the same element. Potassium can exist as three different atoms. All three atoms contain 19 protons, but one potassium atom has 20 neutrons, another 21 neutrons and yet another has 22 neutrons. • Atoms that have the same number of protons but differ in their number of neutrons are called isotopes. • As you would expect, if different isotopes have different numbers of neutrons, they will have different masses. The mass number of an atom is the sum of the protons and neutrons found in the nucleus of that atom.

  35. If we look at the potassium isotopes above, the isotope containing 19 protons and 20 neutrons will have a mass number of ___________. We call this isotope potassium-39. • The isotope that has 19 protons and 21 neutrons will have a mass number of ___________ and is called potassium-40. • Chemists have designed a symbol for each isotope that includes the element’s symbol, its atomic number (Z) and its mass number (A). 19 + 20 = 39 19 + 21 = 40 • The symbol for potassium-39 would be: • The symbol for potassium-40 would be:

  36. Atomic Mass The masses of individual atoms are expressed as atomic mass units (amu) or µ. The atomic mass unit is defined as 1/12 the mass of a carbon-12 atom. This means a proton or a neutron has mass equal to approximately one atomic mass unit. In many cases the amount of each isotope in the sample, or its relative abundance, can be determined using a mass spectrometer. The relative abundance of an isotope is the percent of each isotope found in an average sample of the element. You have noticed that the atomic mass shown for each element on a periodic table is rarely a whole number. This is because it is actually an average mass of all isotopes of that element.

  37. How to Calculate Average Atomic Mass To determine the average atomic mass, you first need to determine what the mass contribution is for the isotope. mass contribution = (mass)(relative abundance) Once all the mass contributions have been determined, you simply add up the numbers to find your average atomic mass. • Example: • Find the atomic mass of magnesium, using the information provided.

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